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Chem C1403 Lecture 9 Wednesday, October 5, 2005

Chem C1403 Lecture 9 Wednesday, October 5, 2005. Today we’ll revisit the discharge lamp experiments with an atomic and electronic interpretation based on the Bohr atom.

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Chem C1403 Lecture 9 Wednesday, October 5, 2005

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  1. Chem C1403 Lecture 9 Wednesday, October 5, 2005 Today we’ll revisit the discharge lamp experiments with an atomic and electronic interpretation based on the Bohr atom. We’ll also review some of the key equations of Einstein, Bohr and deBroglie that provide insight to the paradigm shifts that lead to modern quantum mechanics. We’ll then begin an examination of the modern quantum mechanical interpretation of the H atom.

  2. Robert Grubbs: Nobel Prize in Chemistry: 2005 New methods of forming polymers. Former graduate student at Columbia and former football and basketball (he’s 6’ 5” tall) opponent!

  3. Maxwell: Light consists of waves (energy is propagated by waves): Energy is spread over space like an oscillating liquid. Maxwell’s theory is called the classical theory of light. Key equations: c =   (Gk lambda), (Gk nu) c = speed of light wave wave propagation  = wavelength,  = frequency James Clerk Maxwell 1831-1879 Classical Paradigm: Energy carried by a light wave is proportional to the Amplitude of wave. Big wave, small wave.  Low Frequency l   High Frequency

  4. Waves and light c = = 3.0 x 108 m-s-1 = 3.0 x 1017 nm-s-1 = c/,  = c/ c = speed of light  = frequency of light = speed of light A computation: What is the frequency of 500 nm light? Answer: = c/ = (3.0 x 1017 nm-s-1)/  = 6 x 1014 s-1

  5. Fig 16-5

  6. 4 paradoxes that doomed the classical paradigm of light (and matter) Death spiral of the electron Ultraviolet catastrophe Photoelectric effect Line spectra of atom

  7. Planck explains the ultraviolet catastrophe by quantizing the energy of light. Light can only have energies given by E = hThe value of h =6.6 x 10-34 Js fits experiment! if E can be anything Max Planck Nobel Prize 1918 “for his explanation of the ultraviolet catastrophe”, namely E = h, the energy of light is bundled and comes in quanta. if E = h

  8. Einstein’s explanation of the photoelectric effect” Light consists of photons which carry quanta of energy.

  9. Blue light kicks out Electrons even at very low amplitude! Red light is “inert” to kicking out electrons, no matter what the amplitude of the light! E2 - E1 = h Albert Einstein Nobel Prize 1921 “For his explanation of the photoelectric effect”, namely, E2 - E1 = h, light is quantized as photons. The slope of KEMax vs  is h!!!!!

  10. = KEMAX = h - h0 (Excess kinetic energy of the electron) (work function to remove the electron from metal) Key equation: KEMAX = h( - 0) Slope = h!!!!

  11. Interpreting data: Which metal takes least energy to eject an electron?

  12. The predicted death spiral of the Rutherford atom. The Rutherford atom. Bohr solved this paradox and the paradox of the line spectra of atoms with an assumption and some algebra

  13. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/bohr.htmlhttp://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/bohr.html Wavelength (l) Color 656.2 red 486.1 blue-green 434.0 blue-violet 410.1 violet

  14. Lamp left Lamp right You’ll see something like this on the front podium. You’ll see something like this through your diffraction glasses! Let’s do an experiment: Look at the discharge lamps through the diffraction glasses. They work just like a prism and break up light into its components. Notice the dark spots between the “lines” of the different colors. The number and positions of the lines are the unique signature of the elements. A lab experiment. Note the number and color of the lines. See if you can identify the element.

  15. Unknown

  16. Certain orbits have special values of angular momentum and do not radiate: mever = n(h/2) n = 1, 2, 3,….infinity (This solves the death spiral problem) The energy and frequency of light emitted or absorbed is given by the difference between the two orbit energies, e.g.,E(photon) = E2 - E1 (Energy difference) = h (This solves the line spectrum paradox) Niels Bohr Nobel Prize 1922 “the structure of atoms and the radiation emanating from them” The basis of all photochemistry and spectroscopy!

  17. But there was more, much more that Bohr did than qualitatively take care of the two remaining paradoxes. He then applies some quantitative thinking to figure out what the size of the H atom was based on his hypothesis and then to compute the energies of the jumps between orbits!

  18. By solving the line spectrum paradox, the Bohr model allowed the computation of the energy of an electron in a one electron atom: En = -Ry(Z2/n2) Ry = 2.18 x 10-18 J The results of his computations compared very favorably with experimental data for one electron atoms, but failed completely for atoms with more than one electron! Something was still missing!

  19. What was missing? The electron was being treated as a particle. If waves can mimic particles, then perhaps particles can mimic waves. Light: E = h (Planck) Mass: E = mc2 (Einstein) then h  h(c/)=mc2 (de Broglie) Light = Matter Louis de Broglie 1892-1987 Nobel Prize 1929 “for his discovery of the wave nature of electrons”  = h/mv Two seemingly incompatible conceptions can each represent an aspect of the truth ... They may serve in turn to represent the facts without ever entering into direct conflict. de Broglie, Dialectica

  20. Traveling waves and standing waves Every wave has a corresponding “wavefunction” that completely describes all of its properties.  A circular standing wave With 7 wavelengths around the circle. Localized in space (on an atom!) Light as a traveling wave. No beginning and no end

  21. Wavy cows? = h/mv The value of h = 6.6 x 10-34 Js Electrons show wave properties, cows do not. The wave properties of matter are only apparent for very small masses of matter.

  22. A computations of the wavelength of a macroscopic object (smaller than a cow): A baseball of 0.145 kg of mass, traveling at 30 m-s-1 DeBrolie equation: = h/mv h = 6.63 x 10-34 J-s m = 0.145 kg, v = 30 m-s-1 = h/mv = 6.63 x 10-34 J-s/(0.145 kg, v = 30 m-s-1)  = 1.5 x 10-34 m = 1.5 x 10-24 Å This is such a small number that it cannot be measured and completely masks the wave behavior of macroscopic objects.

  23. Wave, particles and the Schroedinger equation Constructive interference Diffraction patterns: Constructive and destructive interference, the signature characteristic of waves. Destructive interference

  24. The electron as a bound wave: what is its wavefunction? Schroedinger: wave equation and wavefunctions

  25. Wavefunctions and orbitals An orrbital is a wavefunction Obital: defined by the quantum numbers n, l and ml (which are solutions of the wave equation) Orbital is a region of space occupied by an electron Orbitals has energies, shapes and orientation in space p orbitals s orbitals

  26. Sizes, Shapes, and orientations of orbitals n determines size; l determines shape ml determines orientation

  27. The hydrogen s orbitals (solutions to the Schroedinger equation. Fig 16-19

  28. Electron probability x space occupied as a function of distance from the nucleus

  29. The p orbitals of a one electron atom px py pz

  30. The d orbitals of a one electron atom Fig 16-21

  31. The f orbitals of a one electron atom

  32. Quantum Numbers (QN) Principal QN: n = 1, 2, 3, 4…… Angular momentum QN: l = 0, 1, 2, 3…. (n -1) Rule: l = (n - 1) Magnetic QN: ml = …-2, -1, 0, 1, 2, .. Rule: -l….0….+l Shorthand notation for orbitals Rule: l = 0, s orbital; l = 1, p orbital; l = 2, d orbital l = 3, f orbital 1s, 2s, 2p, 3s, 3p, 4s, 4p, 4d, etc.

  33. The energy of an orbital of a hydrogen atom or any one electron atom only depends on the value of n shell = all orbitals with the same value of n subshell = all orbitals with the same value of n and l an orbital is fully defined by three quantum numbers, n, l, and ml Each shell of QN = n contains n subshells n = 1, one subshell n= 2, two subshells, etc Each subshell of QN = l, contains 2l + 1 orbitals l = 0, 2(0) + 1 = 1 l = 1, 2(1) + 1 = 3

  34. Nodes in orbitals: 2p orbitals: angular node that passes through the nucleus Orbital is “dumb bell” shaped Important: the + and - that is shown for a p orbital refers to the mathematical sign of the wavefunction, not electric charge!

  35. Nodes in orbitals: 3d orbitals: two angular nodes that passes through the nucleus Orbital is “four leaf clover” shaped d orbitals are important for metals

  36. The fourth quantum number: Electron Spin ms = +1/2 (spin up) or -1/2 (spin down) Spin is a fundamental property of electrons, like its charge and mass. (spin up) (spin down)

  37. A singlet state A triplet state

  38. Electrons in an orbital must have different values of ms This statement demands that if there are two electrons in an orbital one must have ms = +1/2 (spin up) and the other must have ms = -1/2 (spin down) This is the Pauli Exclusion Principle An emptyorbital is fully described by the three quantum numbers: n, l and ml An electron in an orbital is fully described by the four quantum numbers: n, l, mland ms

  39. Summary of quantum numbers and their interpretation

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